1. Acids and Bases: (An Introduction) Chemistry 12◊ Chapter 14

2. Properties of Acids: Turn blue litmus paper red Neutralize the properties of bases React with certain metals to produce hydrogen gas React with carbonate compounds to produce carbon dioxide gas Have a sour taste Are electrolytes Have a pH less than 7

3. Common Acids: Sulfuric AcidH 2 SO 4 Nitric AcidHNO 3 Phosphoric AcidH 3 PO 4 Hydrochloric AcidHCl Acetic AcidCH 3 COOH Carbonic Acid H 2 CO 3 Battery acid Used to make fertilizers and explosives Food flavoring Stomach acid Vinegar Carbonated water

4. Properties of Bases: Turn red litmus paper blue Turn the indicator phenolphthalein from colorless to red Neutralize the properties of acids Have a bitter taste Are electrolytes Are slippery to the touch Have a pH greater than 7

5. Common Bases: Sodium hydroxideNaOHlye (caustic soda) Potassium hydroxideKOHlye (caustic potash) Magnesium hydroxideMg(OH) 2 milk of magnesia Calcium hydroxideCa(OH) 2 slaked lime Ammonia waterNH 3 H 2 Ohousehold ammonia.

6. Definition of Acids and Bases: An operational definition is a definition based on observed experimental properties. Example: An acid is that it is a substance that turns blue litmus paper red and has a pH less than 7. A conceptual definition attempts to explain why a substance behaves the way it does. We will be looking at a couple of conceptual definitions of acids For example: Arrhenius Theory and Bronsted-Lowery Acid Base Theory

7. 1. Arrhenius Theory: An acid is a substance that produces H + ions in solution (H + always combine with at least one water molecule to produce hydronium ions, H 3 O +. HCl(g) + H 2 O(l) -- > H 3 O + (aq) + Cl - (aq) A base is a substance that produces OH - ions in solution. NaOH(s) + H 2 O(l) → Na + (aq) + OH-(aq)

8. Limitations of Arrhenius Theory: Acids (like HCl) are able to be neutralized by NaOH to produce water: But what about a base like NH 3 ….. Not all bases have OH- ions, NH 3 in H 2 O does but, if NH 3 reacts as a gas with HCl gas, NH 4 Cl is still formed, but is not a solution, so this theory won’t always work!

9. 2. Bronsted-Lowry Theory: An acid is a proton (H + ion) donor. A base is a proton (H + ion) acceptor. H 3 O + is a hydronium ion (aka. THE PROTON) HCl is the acid (aka the donor) H 2 O is the base (aka the acceptor)

10. Conjugate Acid-Base Pairs: Conjugate = “linked together” A conjugate acid is the substance that has accepted the proton (gained an H + ) A conjugate base is the substance that has lost the proton. Each acid has a conjugate base and each base has a conjugate acid.

11. Conjugate Acid-Base Pairs:

12. How many protons can be lost? Bronsted-Lowry acids can be: Monoprotic acids are capable of losing one proton: HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) HNO 3(aq) + H 2 O (l)  H 3 O + (aq) + NO 3 - (aq) Diprotic acids are capable of losing two protons (in more then one step): H 2 SO 4(aq) + H 2 O (l)  HSO 4 - (aq) + H 3 O + (aq) HSO 4 - (aq) + H 2 O (l)  SO 4 2- (aq) + H 3 O + (aq)

13. Amphoteric: Substances that can act like an acid in one reaction and like a base in another type of reaction. Example: hydrogen carbonate ( HCO 3 -) 1. HCO 3 - + OH- CO 3 -2 + H 2 O (donates a H +, so acts like an acid) 2. HCO 3 - + H 3 O+ H 2 CO 3 + H 2 O (accepts a H +, so acts like a base)

14. 3. Lewis Theory: Bases donate pairs of electrons and acids accept pairs of electrons. (Acid/base reaction is the donation of an electron pair to create a new covalent bond ) A Lewis acid is any substance, such as the H + ion, that can accept a pair of non-bonding (lone) electrons. A Lewis acid is an electron-pair acceptor. A Lewis base is any substance, such as the OH - ion, that can donate a pair of non-bonding electrons. A Lewis base is an electron-pair donor.

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